Chapter 4. Lecture Presentation

Transcription

1 Chapter 4 Lecture Presentation

2 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

3 Goal : Given its name, write its correct symbol; from the symbol. Write the correct name.

4 Elements: are pure substances from which all other things are built. are listed on the periodic table.

5 Element names come from planets, mythological figures, minerals, colors, geographic locations, and famous people.

6 Chemical symbols represent the names of the elements. consist of one to two letters and start with a capital letter. Carbon, C One-Letter Symbols Two-Letter Symbols C carbon Co cobalt N nitrogen Ca calcium F fluorine Al aluminum O oxygen Mg magnesium

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8 Write the correct chemical symbols for each of the following elements: A. iodine B. iron C. magnesium D. zinc E. nitrogen

9 Give the names of the elements with the following symbols: A. P B. Al C. Mn D. H E. K

10 Mercury (Hg) Is a silvery, shiny element that is liquid at room temperature. Can enter the body by: mercury vapor inhalation contact with the skin ingestion of water or food contaminated with mercury Once mercury has entered the body, it destroys proteins, disrupts cell function. Long-term exposure can Damage the brain and kidneys Cause mental retardation Decrease physical development

11 Mercury contamination comes from industrial wastes. fish and seafood. batteries. compact fluorescent bulbs. Fish absorb mercury. Big fish eat lots of small fish, end up with more mercury.

12 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

13 Goal : Use the periodic table to identify the group and the period of an element; identify the element as a metal, a nonmetal, or a metalloid.

14 The Periodic Table The periodic table organizes 118 elements into groups with similar properties and places them in order of increasing atomic mass.

15 Groups and Periods In the periodic table, elements are arranged according to properties. Vertical columns represent groups of elements and horizontal rows represent periods.

16 Groups and Periods Group numbers are written at the top of each vertical column (1-18). Periods are numbered 1-7

17 Metals, Nonmetals, and Metalloids The heavy zigzag line separates metals and nonmetals. Metals are located to the left. Nonmetals are located to the right. Metalloids are located along the heavy zigzag line.

18 Metals Solid at room temp Exception: mercury Shiny! Ductile (shaped into wires) Malleable (hammered flat into sheets) Good conductors High melting points Silver (Metal) Antimony (metalloid) Sulfur (nonmetal) Nonmetals Dull (not shiny ) brittle Not ductile Not malleable Poor conductors Low melting points Low densities (many are gasses at room temp.) Metalloids have a combination of metal and nonmetal properties.

19 Characteristics of Metalloids Metalloids, located along the heavy zigzag line on the periodic table exhibit properties of metals and nonmetals. are better conductors than nonmetals but not as good as metals. are used as semiconductors and insulators, because they can be modified to function as conductors or insulators. Silver (Metal) Antimony (metalloid) Sulfur (nonmetal)

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21 Study Check Identify each of the following elements as a metal, a nonmetal, or a metalloid: A. sodium B. chlorine C. silicon D. iron E. carbon

22 Study Check List all of the elements that match the description. A. metals in Group 14 Sn, Pb, C, Si, Ge B. nonmetals in Group 15 Bi, N, P, As, Sb C. metalloids in Group 14 C, Si, Ge, Sn, Pb

23 Group Names

24 Alkali Metals Group 1, the alkali metals, includes the following: lithium (Li) sodium (Na) potassium (K) rubidium (Rb) cesium (Cs) Francium (Fr) Soft Shiny Good conductors Low melting points React vigorously with water!

25 Alkaline Earth Metals Group 2 elements, the alkaline earth metals, are shiny but not as reactive as Group 1A metals. They include the following: beryllium (Be) magnesium (Mg) calcium (Ca) strontium (Sr) barium (Ba) radium (Ra) Strontium gives the red color in fireworks.

26 Halogens Group 17, the halogens, includes the following: fluorine (F) chlorine (Cl) bromine (Br) iodine (I) astatine (At)

27 Noble Gases Group 18, the noble gases, include: Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) Xenon (Xe) Radon (Rn)

28 Study Check Identify the element described by each of the following groups and periods: 1. Group 17, Period 4 2. Group 2, Period 3 3. Group 15, Period 2

29 FYI: Elements Essential to Health Of all the elements, 23 are essential for the well-being and survival of the human body.

30 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

31 Goal : Describe the electrical charge and location in an atom for a proton, a neutron, and an electron

32 The Atom An atom is the smallest particle of an element that retains the characteristics of that element. Aluminum foil contains atoms of aluminum.

33 Atoms are made of subatomic particles Atoms contain the following subatomic particles: Protons Neutrons Electrons

34 Atoms are made of subatomic particles Atoms contain the following subatomic particles: protons have a positive (+) charge neutrons have no charge (neutral) electrons have a negative (-) charge

35 Structure of the Atom An atom consists of a nucleus, located in the center of the atom, that contains protons and neutrons and represents most of the mass of an atom. electrons that occupy a large, empty space around the nucleus. Protons positive charge Neutrons neutral Electrons negative charge

36 Structure of the Atom In an atom, the protons and neutrons that make up almost all the mass are packed into the tiny volume of the nucleus. The rapidly moving electrons (negative charge) surround the nucleus and account for the large volume of the atom.

37 Dalton s Atomic Theory In Dalton s atomic theory, atoms are tiny particles of matter. of an element are similar to each other and different from those of other elements. of two or more different elements combine to form compounds. are rearranged to form new combinations in a chemical reaction. Atoms are never created or destroyed during a chemical reaction.

38 Mass of the Atom Because the mass of subatomic particles are so small, chemists use a very small unit of mass called the atomic mass unit (amu). 1 amu has a mass equal to 1/12 of the mass of the carbon-12 atom that contains six protons and six neutrons. 1 amu = 1 Dalton (Da) in biology. 1 amu = 1.66 x kg Electrons have such a small mass that they are not included in the mass of an atom.

39 If a proton has a mass of 1.67 x kg, what is its mass in amu?

40 Subatomic Particles in the Atom Protons and neutrons have a very small mass. Electrons are 1800 times smaller than protons and neutrons.

41 Study Check Which of the following subatomic particles fits each of the descriptions below? protons, neutrons, or electrons A. found outside the nucleus B. have a positive charge C. have mass but no charge

42 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

43 Goal : Given the atomic number and the mass number of an atom, state the number of protons, neutrons, and electrons.

44 Atomic Number Each element is assigned an atomic number. The atomic number is equal to the number of protons in an atom (and is always a whole number). The same for all atoms of an element Appears above the symbol of an element in the periodic table.

45 Atomic Number - examples Atomic number = number of protons the atomic number of H is 1; every H atom has one proton. the atomic number of C is 6; every C atom has six protons. the atomic number of Cu is 29; every Cu atom has 29 protons.

46 Atomic Number - example All atoms of lithium (left) contain three protons, and all atoms of carbon (right) contain six protons.

47 Atoms are Neutral For neutral atoms, the net charge is zero. number of protons = number of electrons Aluminum has 13 protons and 13 electrons. The net (overall) charge is zero. 13 protons (13+) + 13 electrons (13 ) = 0

48 Study Check Use the periodic table to fill in the atomic number, number of protons, and number of electrons for each of the following elements: Element N Zn S Atomic Number Protons Electrons

49 Mass Number The mass number represents the number of particles in the nucleus. is equal to the number of protons + the number of neutrons. is always a whole number. does not appear in the periodic table.

50 Study Tips Number of protons = atomic number Number of protons + neutrons = mass number Number of neutrons = mass number atomic number Charge of atom = Number of electrons + Number of protons Number of electrons = charge number of protons

51 Study Check An atom of lead (Pb) has a mass number of 207. A. How many protons are in the nucleus? B. How many neutrons are in the nucleus? C. How many electrons are in the atom?

52 Study Check An atom of titanium (Ti) has a mass number of 44. A. How many protons are in the nucleus? B. How many neutrons are in the nucleus? C. How many electrons are in the atom?

53 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

54 Goal : Determine the number or protons, electrons, and neutrons in one or more of the isotopes of an element; calculate the atomic mass of an element using the percent abundance and mass of its naturally occurring isotopes.

55 Isotopes Isotopes are atoms of the same element. have the same number of protons but different numbers of neutrons. (different mass numbers) can be distinguished by their atomic symbols.

56 Atomic Symbols: Subatomic Particles Given the atomic symbols, determine the number of protons, neutrons, and electrons. 16 O P Zn 15 30

57 Study Check Naturally occurring carbon consists of three isotopes: 12 C, 13 C, and 14 C. State the number of protons, neutrons, and electrons in each of the three isotopes.

58 Study Check Write the atomic symbols for atoms with the following subatomic particles: A. 8 protons 8 neutrons 8 electrons B. 17 protons 20 neutrons 17 electrons C. 47 protons 60 neutrons 47 electrons

59 Study Check 1. Which of the pairs below are isotopes of the same element? 2. Which of the pairs below have the same number of neutrons? A B C 15 X 15 X 12 X 14 X 15 X 16 X

60 Naturally Occurring Isotopes Most elements have several isotopes that occur in nature (vs. made in a lab.) These isotopes are called naturally occurring. Chlorine has 3 isotopes: 35 Cl 36 Cl 37 Cl Only 35 Cl and 37 Cl happen naturally. 36 Cl has to be made in a lab. So chlorine has 2 naturally occurring isotopes.

61 Atomic Mass Atomic mass is the weighted average of all naturally occurring isotopes of that element. number on the periodic table below the chemical symbol. Chlorine, with two naturally occurring isotopes, has an atomic mass of amu.

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63 Isotopes of Magnesium Magnesium, with three naturally occurring isotopes, has an atomic mass of amu.

64 Isotopes of Magnesium Magnesium s atomic mass: amu.

65 Atomic Mass of Some Elements

66 Calculating Atomic Mass To calculate atomic mass, use the experimental percent abundance of each isotope of the element. multiply the percent abundance (divided by 100) by the atomic mass of that isotope. sum the total mass of all isotopes.

67 Calculating Atomic Mass

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69 Study Check Gallium is an element found in lasers used in CD players. In a sample of gallium, there is 60.10% of 69 Ga atoms (atomic mass ) 39.90% of 71 Ga atoms (atomic mass ) What is the atomic mass of gallium?

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71 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

72 Goal: Describe the energy levels, sublevels, and orbitals for the electrons in an atom.

73 Electromagnetic Radiation We experience electromagnetic radiation in different forms, such as light, the colors of a rainbow, or X-rays.

74 Electromagnetic Radiation Light and other electromagnetic radiation consists of energy particles that move as waves of energy. The distance between the peaks of waves is called the wavelength. High-energy radiation has shorter wavelengths. Low-energy radiation has longer wavelengths. All electromagnetic radiation travels at the speed of light! (3 x 10 8 m/s!) Light is pure energy!

75 Electromagnetic Spectrum The Electromagnetic Spectrum above shows all types electromagnetic waves. Many of which we use every day!

76 A rainbow appears when sunlight reflects through a raindrop and is separated into its different EM radiation. The same thing happens when you shine sunlight through a glass prism!

77 Atomic Spectrum Scientists found that when they used a different light source than the sun, they didn t always get a full rainbow! They found that each element had it s own fingerprint based on which light it produced. We now call it an element s atomic spectrum

78 Bohr model and bookcase example The lines of color in an atomic spectrum are caused by the behavior of that element s electrons. Electrons occupy specific areas around the atom they are a part of. Each electron has a specific energy level which corresponds to the ring the electron resides in.

79 Electrons and Energy Levels Electrons with the same energy are grouped in the same energy level. Energy levels are assigned values called principal quantum numbers (n), (n = 1, n = 2, ). An electron can have only the energy of one of the energy levels in an atom.

80 Electron Energy Levels Bohr model In an atom, each electron has a specific energy, known as its energy level, which is assigned principal quantum numbers (n) = (n = 1, n = 2, ). increases in energy as the value of n increases and electrons are farther away from the nucleus. The energy of an electron is quantized electrons can have only specific energy values.

81 Changes in Electron Energy Level Electrons move to a higher energy level when they absorb energy. When electrons fall back to a lower energy level, light is emitted. The energy emitted or absorbed is equal to the differences between the two energy levels.

82 Sublevels Bookshelf with sublevels It is the arrangement of electrons that determines the physical and chemical properties of an element. Each energy level consists of one or more sublevels. The number of sublevels in an energy level is equal to the principal quantum number n of that energy level. The sublevels are identified as s, p, d, and f. The order of sublevels in an energy level is s < p < d < f

83 Sublevels Up to 2 electrons can fit in each box (orbital)

84 s Orbitals The location of an electron is described in terms of probability. Orbitals are a threedimensional volume in which electrons have the highest probability of being found. The s orbitals are spheres. (a) The electron cloud of an s orbital represents the highest probability of finding an s electron. (b) The s orbitals are shown as spheres. The sizes of the s orbitals increase because they contain electrons at higher energy levels.

85 p Orbitals There are three p orbitals, starting with n = 2. Each p orbital has two lobes, like a balloon tied in the middle, and can hold a maximum of two electrons. The three p orbitals are arranged perpendicular to each other along the x, y, and z axes around the nucleus.

86 p Orbitals A p orbital has two regions of high probability, which gives a dumbbell shape. (a) Each p orbital is aligned along a different axis from the other p orbitals. (b) All three p orbitals are shown around the nucleus.

87 d Orbitals There are five d orbitals, starting with n = 3. Four of the five d orbital has four lobes, in the shape of a 4-leaf clover. The only difference between them is their location. Shape is identical. The 5 th d orbital looks like a p orbital with a donut around its middle.

88 d Orbitals Each of the d sublevels contains five d orbitals. Four of the five d orbitals consist of four lobes that are aligned along or between different axes. One d orbital consists of two lobes and a doughnutshaped ring around its center.

89 Sublevels Up to 2 electrons can fit in each box (orbital)

90 Each orbital can hold to two electrons.

91 Orbital Capacity and Electron Spin electrons in the same orbital repel each other. electrons in the same orbital must have their magnetic spins cancel (they must spin in opposite directions). We can represent magnetic spins with arrows Summary: Each orbital can have up to 2 electrons with one spin up and the other spin down

92 The Pauli exclusion principle states that each orbital can hold a maximum of two electrons. electrons in the same orbital repel each other. electrons in the same orbital must have their magnetic spins cancel (they must spin in opposite directions).

93 Number of Electrons in Sublevels There is a maximum number of electrons that can fill each sublevel. Each s sublevel has one orbital and can hold a maximum of two electrons. Each p sublevel has three orbitals and can hold a maximum of six electrons. Each d sublevel has five orbitals and can hold a maximum of 10 electrons. Each f sublevel can has 7 orbitals and can hold a maximum of 14 electrons.

94 Number of Electrons in Sublevels

95 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

96 Goal: Draw the orbital diagram and write the electron configuration for an element.

97 An orbital diagram shows the placement of electrons in the orbitals in order of increasing energy. This is much easier than trying to draw all the orbital shapes. There is 1 s orbital, so there is 1 box. There are 3 p orbitals, so there are 3 boxes representing it. Etc.

98 Orbital Diagrams H and He We fill an orbital diagram with the number of electrons an atom has. Element Atomic # # of electrons Hydrogen (H) Helium (He) Begin at the bottom of the chart, putting two electrons in each orbital (box). Use an arrow to represent an electron. Hydrogen Helium If there are two electrons in one orbital, make one up and one down

99 Orbital diagrams can also be written on one line to save space. They are written in the order of lowest energy to highest. Hydrogen Helium

100 Finally, orbital diagrams can also be written using the orbital names and number of electrons as: We call this format the electron configuration of that atom or element.

101 Electron Configurations Chemists use a notation called electron configuration to indicate the placement of electrons in an atom. show how electrons fill energy levels and sublevels in order of increasing energy. write an abbreviated form using a noble gas to represent all electrons preceding it. Electron Configuration for Carbon

102 Period 2: Li Element Atomic # # of electrons Lithium

103 Period 2: Be Element Atomic # # of electrons Beryllium

104 Period 2: B, C, and N Element Atomic # # of electrons Boron Carbon Nitrogen Place one electron in each p orbital before doubling up. (same for d and f orbitals)

105 Period 2: O, F, and Ne Element Atomic # # of electrons Oxygen Fluorine Neon Place one electron in each p orbital before doubling up. (same for d and f orbitals)

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107 Period 3: Sodium to Argon

108 Electron Configurations and the Periodic Table The electron configurations of elements are related to their positions on the periodic table. Different sections or blocks correspond to sublevels s, p, d, and f.

109 Blocks on the Periodic Table 1. The s block contains elements in 1 and 2. This means the final one or two electrons are in the s sublevel. 2. The p block consists of elements in Group 13 to Group 18. There are six p block elements in each period, because three p orbitals can hold a maximum of six electrons.

110 Blocks on the Periodic Table 3. The d block, which contains transition elements, first appears after calcium (atomic number 20). There are 10 elements in the d block, because five d orbitals can hold a maximum of 10 electrons. 4. The f block, the inner transition elements, is the two rows of elements at the bottom of the periodic table. There are 14 elements in each f block, because seven f orbitals can hold a maximum of 14 electrons.

111 Writing Configurations Using Sublevel Blocks STEP 1 Locate the element on the periodic table. STEP 2 Write the filled sublevels in order, going across each period. Example: Chlorine

112 Electron Configurations: d and f block Beginning in Period 4, the 4s sublevel fills before the 3d sublevel, because the 3d sublevel is slightly lower in energy than the 4s sublevel. the 5s sublevel fills before the 4d sublevel. the 6s sublevel fills before the 5d sublevel.

113 Study Check Use the sublevel blocks on the periodic table to write the electron configuration for selenium.

114 d block Sublevel Exceptions For chromium (Cr), moving one of the 4s electrons to the 3d sublevel adds stability with a half-filled d subshell, and the resulting configuration is 4s 1 3d 5. For copper (Cu), moving one of the 4s electrons to the 3d sublevel adds stability with a filled d subshell, and the resulting configuration is 4s 1 3d 10.

115 Study Check Use the sublevel blocks on the periodic table to write the electron configuration for rhenium (Re).

116 Study Check Use the periodic table to give the symbol and name for the element with the electron configuration of 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7.

117 Study Check Write the electron arrangement for the following elements: C Si O 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f

118 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

119 Goal: Use the electron configurations of elements to explain the trends in periodic properties.

120 Valence Electrons Most properties of elements are due to the behavior of the electrons in the element s outermost orbitals. Valence electrons are the number of electrons in the outermost energy level.

121 Valence Electrons

122 Study Check Using the periodic table, how many valence electrons do each element have? A. calcium B. lead

123 Lewis Symbols Lewis symbols represent the valence electrons as dots placed on sides of the symbol for an element. One to four valence electrons are arranged as single dots. Five to eight valence electrons are arranged with at least one pair of electrons around the symbol for the element. H He B S Kr Remember, with valence electrons, it doesn t matter which row (period) you are in. Only the column (group).

124 Study Check Write the electron-dot (Lewis) symbol for each of the following elements: Cl C N

125 Atomic Size Atomic size is determined by the atom s atomic radius, the distance between the nucleus and the outermost electrons.

126 Atomic Size The atomic size increases going down a group and right to left across a period.

127 Atomic size increases top to bottom down a group Because each period adds more orbitals which are increasingly farther from the nucleus. Atomic size increases from left to right Because as you right along a period, you are adding more protons and electrons (but have the same orbitals.) This creates a stronger attraction between the nucleus (protons) and the electrons and it sucks the electrons in tighter. Making the atom smaller.

128 Study Check Given the elements C, N, and Cl, A. which is the largest atom?

129 Ionization Energy Ionization energy is the energy required to remove one of the outermost (valence) electrons. As the distance from the nucleus to the valence electrons increases, the ionization energy decreases. The ionization energy is low for metals and high for the nonmetals.

130 Ionization Energy Ionization energy increases up a group and increases going across a period from left to right.

131 Study Check Given the elements C, N, and Cl, B. which has the highest ionization energy?

132 Metallic Character An element with metallic character is one that loses valence electrons easily (low ionization energy). Metallic character is more prevalent in metals on the left side of the periodic table. is less for nonmetals on the right side of the periodic table that do not lose electrons easily. decreases going down a group, as electrons are farther away from the nucleus.

133 Metallic Character The metallic character of the s and p block elements increases going down a group and increases going from right to left across a period.

134 Summary : Periodic Trends

135 4.1 Elements and Symbols 4.2 The Periodic Table 4.3 The Atom 4.4 Atomic Number and Mass Number 4.5 Isotopes and Atomic Mass 4.6 Electron Energy Levels 4.7 Electron Configurations 4.8 Trends in Periodic Properties

136 Concept Map